CHAPTER 2: THE CHEMICAL BASIS OF LIFE. 3. Give the chemical symbol for the naturally occurring elements in humans

CHAPTER 2: THE CHEMICAL BASIS OF LIFE OBJECTIVES 1. Define the terms chemistry and matter. 2. Define the terms atom and element. 3. Give the che...
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CHAPTER 2: THE CHEMICAL BASIS OF LIFE

OBJECTIVES 1.

Define the terms chemistry and matter.

2.

Define the terms atom and element.

3.

Give the chemical symbol for the naturally occurring elements in humans.

4.

List the four major elements that compose the human body.

5.

Describe and distinguish between the three subatomic particles in terms of charge, weight, and location of each. Sketch a diagram to illustrate their relationship.

6.

Distinguish between the atomic number and atomic weight of an atom of an element.

7.

Discuss how isotopes of atoms of a particular element differ.

8.

Given the atomic number of an atom, you should be able to determine the following: a. b. c. d. e.

the number of protons; the number of electrons; the electron configuration of the atom; the number of valence electrons; how that atom will react.

9.

Explain how atoms react with one another.

10.

Distinguish between ionic, covalent, and hydrogen bonds, and give an example of a molecule (or macromolecule) that demonstrates each.

11.

Name the three types of chemical reactions.

12.

Compare and contrast the major divisions (types of chemical reactions) of metabolism, in terms of a general descriptive sentence, additional descriptive terms, how energy is involved, whether bonds are formed or broken, and how water is involved. Also write a chemical reaction for each and give an example important in human metabolism.

13.

List the four factors that affect the rate at which chemical reactions occur.

14

CHAPTER 2: CHEMISTRY OBJECTIVES 14.

Distinguish between organic and inorganic compounds.

15.

List five inorganic substances of importance to humans.

16.

Discuss the unique structure of a water molecule and name the bonds that hold liquid water together.

17.

List and discuss the characteristics of water.

18.

List the major electrolytes released by inorganic salts when placed in water and explain how these electrolytes are needed for metabolic reactions.

19.

Describe what happens to an acid and base when they are placed in water, and discuss the significance of these products in the human body.

20.

Illustrate the pH scale, denoting acid, neutral, and basic (alkaline) pH values. Also denote the relationship between [H+] to [OH-] at each of the above pH’s, and show approximately where on that scale the following substances would fall: acetic acid, distilled water, blood and ammonia.

21.

Using the scale above, plot the pH values of the compounds you tested in lab.

22.

Name the value of physiological pH.

23.

Define the term buffer, and explain how the carbonic acid buffering system works in humans.

24.

List the four major organic substances needed for human survival (i.e. proteins, carbohydrates, lipids and nucleic acids), name the building blocks that compose each, and give a general function for each.

25.

Name the three types of atoms that compose sugars and lipids.

26.

Name three monosaccharides and three disaccharides.

27.

Name two polysaccharides, indicate whether each is a plant or animal carbohydrate, and name the tissue where the animal carbohydrate is stored.

28.

Distinguish between the three types of lipids, in terms of structure and function.

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CHAPTER 2: CHEMISTRY OBJECTIVES 29.

Compare and contrast saturated and unsaturated fats.

30.

Name the bond that is formed when two amino acids are joined.

31.

Describe the levels of structural organization of a protein and explain the significance of a protein’s conformation on its overall function.

32.

Define the term denaturation and explain what conditions may cause a protein to become denatured.

33.

List and discuss the many functions of proteins (Which is the most important?)

34.

Discuss the structure of a nucleotide.

35.

Name the type of chemical bond that holds the chains of a DNA molecule together.

36.

List three differences between DNA and RNA.

37.

Name the two types of nucleic acids, describe the structure of each, and give a general function for each molecule.

16

CHAPTER 2: CHEMISTRY I.

II.

DEFINITIONS A.

Chemistry = the study of matter.

B.

Matter = anything that occupies space and has mass; (i.e. solids, liquids, gases)

Structure of Matter A.

Atom =

the smallest particle of an element; the least complex level of organization.

B.

Element =

a basic chemical substance composed of atoms.

C.

1.

Elements are represented by a 1 or 2 letter symbol that are shown in the Periodic Table of the Elements in Appendix A, page 986;

2.

120 elements exist in nature, however only approximately 26 are naturally occurring in humans.

3.

Learn the elements (and their chemical symbol) listed in Table 2.2, page 38.

4.

The most abundant of the naturally occurring elements are carbon (C), Hydrogen (H), Oxygen (O) and Nitrogen (N) = CHON;

Composition of Atoms = 3 Subatomic Particles 1.

Proton =

2.

Neutron =

3.

Electron =

See Fig 2.1 & Table 2.1, p 38.

a positively charged particle in the nucleus of an atom; Mass = 1. an electrically neutral particle in the nucleus of an atom; Mass = 1. an electrically negative particle that revolves around the nucleus; Mass = 0.

SUBATOMIC PARTICLE SUMMARY TABLE (Key on page 36) SUBATOMIC PARTICLE

CHARGE

LOCATION

MASS (WEIGHT)

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CHAPTER 2: CHEMISTRY II.

Structure of Matter D.

Atoms are neutral in charge 1.

E.

The Atomic Number (A#) of an atom represents the number of protons in its nucleus. 1. 2. 3.

F.

The number of protons is equal to the number of electrons.

A# of H = 1 A# of He = 2 A# of O = 8.

The Atomic Weight (AW) of an atom is equal to the number of protons plus the number of neutrons in its nucleus.

1.

Isotopes = atoms of an element that have the same A#’s but different AW’s (i.e. same # of protons, different # of neutrons). a.

The nucleus of some isotopes are stable;

b.

The nucleus of other isotopes are unstable and break apart to become more stable; m

When the nucleus of an atom breaks apart, it releases radioactive energy;

m

Radioactive isotopes have many biological uses. See Clinical Application 2.1, page 40 & 41.

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CHAPTER 2: CHEMISTRY II.

Structure of Matter G.

Electron Configuration 1.

The electrons of an atom are arranged in orbits, shells, or energy levels around the central nucleus; See Fig 2.2 page 40.

2.

A characteristic number of electrons fill each shell: a. b. c.

2 electrons fill the first shell (closest to nucleus); 8 electrons fill the second shell; 8 electrons fill the third shell.

Example 1:

Sodium (Na):Atomic Number = 11; # protons = 11; # electrons = 11.

Example 2:

Sulfur (S):

Atomic Number= 16 # protons= 16 # electrons= 16

19

CHAPTER 2: CHEMISTRY II.

Structure of Matter G.

Electron Configuration 3.

The way in which atoms react with one another (i.e. their chemical properties) is based on the electrons in their outermost shell = VALENCE ELECTRONS a. b. c.

4.

The outermost shell of an atom is called its valence shell. Na has ________ valance electrons; S has _________ valance electrons.

Summary/Overview: Example 1:

Fluorine has an Atomic Number of 9. Draw an atom of fluorine. How and why will fluorine react?

Example 2:

Argon has an Atomic Number of 18. Draw an atom of argon. How and why will argon react?

20

CHAPTER 2: CHEMISTRY III.

CHEMICAL BONDS A.

Atoms form bonds with other atoms to fill their outermost or valence electron shell (energy level). 1.

B.

"Rule of Octets" = except for the first energy level (which contain 2 electrons), atoms react with other atoms so they will have 8 electrons in their valence shell.

Ionic Bonds: See Fig 2.3, page 42. 1. Ions = atoms that have lost or gained electrons to fill their valence shell. a. anion = a negatively ion (Cl-); b. cation = a positively charged ion (Na+). c. An attraction exists between oppositely charged ions and an ionic bond results. Na+ (the anion) donates its outer electron to Cl- (the cation). Table salt or sodium chloride is held together by ionic bonds.

C.

Covalent Bonds: See Fig 2.4, page 42. 1. A covalent bond is formed by the sharing of electrons between atoms. 2. a very strong bond 3. Examples: a. H2 (molecular hydrogen); b. O2 (molecular oxygen); c. H20 (water).

D.

Hydrogen Bonds: See Fig 2.5, page 44. 1. A hydrogen bond is a weak bond formed between hydrogen atoms (that are covalently bonded to another atom) and another atom. 2. Examples include liquid water and DNA chains. 3. These bonds are easily broken and put back together.

CHEMICAL BOND SUMMARY TABLE (Key on page 36) TYPE OF BOND

DEFINITION

DESCRIPTION

EXAMPLE

21

CHAPTER 2: CHEMISTRY IV.

CHEMICAL REACTIONS A.

Definition:

B.

Three Types: 1.

2.

3.

A chemical reaction occurs whenever chemical bonds are formed, rearranged or broken.

Synthesis = the building of a large molecule (polymer) from smaller building blocks (monomers); a.

constructive, anabolic reactions;

b.

Bonds are formed (ENDERGONIC);

c.

Water is usually removed from building blocks to form bond (DEHYDRATION);

d.

Energy A + B ------------> A--B. H20

e.

Example = the building of a large protein (polymer) from many smaller amino acids (monomer).

which

now

hold

chemical

energy

Degradation = breaking a large molecule (polymer) down into its building blocks (monomers); a.

destructive, catabolic, "digestive" reactions;

b.

Bonds are broken releasing chemical energy (EXERGONIC);

c.

Water is used to break bonds (HYDROLYSIS);

d.

H20 A--B -----------------> A + B Energy

e.

Example = digesting a large protein we eat into its amino acid building blocks.

Exchange Reactions involve both degradation and synthesis reactions; a.

A--B + C--D ------> A--C + B--D.

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CHAPTER 2: CHEMISTRY IV.

4.

Chemical Reaction Comparison Table (Key on page 37)

SYNTHESIS REACTIONS

DEGRADATION REACTIONS

GENERAL DESCRIPTION

DESCRIPTIVE TERMS

BOND FORMATION OR BREAKING?

IS ENERGY REQUIRED OR RELEASED? NAME THAT TERM. HOW IS WATER INVOLVED? NAME THAT TERM.

EXAMPLE IN HUMAN METABOLISM

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CHAPTER 2: CHEMISTRY IV.

Chemical Reactions (continued) C.

Factors affecting the rate of chemical reactions: 1.

Particle size: The smaller the particle, the faster the reaction will occur.

2.

Temperature:The higher the temperature, the faster the reaction will occur (up to a point).

3.

Concentration:

4.

Catalysts: a. b.

V.

The greater number of particles in a given space, the faster the reaction.

Definition:

A catalyst is something that speeds the rate of a reaction. Enzymes (special proteins) are biological catalysts.

Types of Compounds in Living Matter A.

Inorganic Compounds are small compounds that do not contain the atoms C and H; Examples include carbon dioxide (CO2) water, salts, acids & bases. 1.

CO2=by-product of cellular respiration.

2.

Water is a polar molecule that demonstrates hydrogen bonding and therefore it possesses very unique characteristics. a.

Water is an excellent solvent (universal?) m m

b.

Many solutes are dissolved in our body’s water. Many ionic compounds (i.e. NaCl) dissociate or break apart in water.

Water participates in many chemical reactions. m Dehydration (synthesis) is when water is removed from adjacent atoms (of molecules) to form a bond between them. m Hydrolysis (degradation) is when water is used to break bonds between molecules.

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CHAPTER 2: CHEMISTRY V.

A.

Inorganic Compounds (continued) 2.

Water (continued) c.

Water is an excellent temperature buffer. m

d.

Water provides an excellent cooling mechanism. m

e.

3.

requires a lot of heat to change from a liquid to a gas; high heat of vaporization.

Water serves as a lubricant m m m

f.

absorbs and releases heat very slowly

mucus; internal organs; joints.

Water is the most abundant component in cells (about 70%).

Salts: See Fig 2.7, page 45. a.

Salts dissociate (ionize) into ions when dissolved in water. m m m

an anion is formed and a cation is formed. Example = NaCl in water. H2O NaCl ---------> Na+ + Cl-

b.

These ions are referred to as electrolytes (charged particles). m m m

Electrolytes must be maintains within a very narrow range in our blood and tissues (i.e. homeostasis); Needed for muscle contraction, nerve impulses, etc.; Examples include Na+, K+, Cl-, Ca+, PO4-; HCO3-, etc.

25

CHAPTER 2: CHEMISTRY V.

A.

Inorganic Compounds (continued) 4/5.

Acids/Bases a.

Acids dissociate (ionize) in water into: m m m

a hydrogen cation, H+, and an anion. Example = HCl (hydrochloric acid). H2O HCl -----------> H+ + Cl-

b.

Bases dissociate (ionize) in water into: m m m

a hydroxyl anion, OH-, and a cation. Example = NaOH (sodium hydroxide). H2O NaOH -----------> Na+ + OH-.

c.

pH Scale: m

See Fig 2.8, page 46.

The relative concentrations of hydrogen ions and hydroxyl ions determine the pH in our blood, fluids, and tissues. pH in body = [H+] + [OH-] .

m

pH = -log[H+];

m

pH Scale ranges from 0 to 14: See Fig 2.8, page 46.

0 ---------------------------7---------------------------14 acid [H+] > [OH-] m

neutral [H+] = [OH-]

basic [H+] < [OH-]

Physiologic pH = 7.4 26

CHAPTER 2: CHEMISTRY V.

A.

4/5.

Acids/Bases (continued) d.

Buffering Systems Definition: m m m m

Buffers are compounds added to solutions to prevent abrupt change in pH.

usually weak acids; function by donating H+ when needed and by accepting H+ when in excess; very important in biological systems! Example = the carbonic acid (H2CO3) buffering system. when pH is rising ----------------> H2CO3 HCO3 7.4 = alkalosis; lethal above 7.8. See Table 2.6, page 48 for Summary of the Inorganic Substances Common in Cells

27

CHAPTER 2: CHEMISTRY VI.

ORGANIC COMPOUNDS A.

General Characteristics: 1.

contain the atoms carbon (and hydrogen);

2.

small molecules (monomers or building blocks) are covalently bonded together to form large polymers or macromolecules;

3.

Water is usually involved in the formation and breakage of bonds between monomers; a. b.

4.

Dehydration Synthesis = removal of water to form a covalent bond between monomers; Hydrolysis = using water to break bonds between monomers.

The four major classes found in cells include: a.

carbohydrates;

b.

lipids;

c.

proteins;

d.

nucleic acids.

28

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) B.

CARBOHYDRATES (sugars) 1.

contain C, H, and O in a 1:2:1 ratio (usually); (Ex: glucose = C6H12O6)

2.

Monomers (building blocks) = monosaccharides; hexoses = simple 6-C sugars; See Fig 2.9, page 48. a. b. c.

3.

glucose, fructose, galactose.

Polymers are formed by dehydration synthesis: a.

disaccharides = 2 monosaccharides covalently bonded together; See Fig 2.10b, page 49. 1. 2. 3.

b.

maltose = glucose + glucose; lactose = glucose + galactose; sucrose = glucose + fructose

polysaccharides = 100 to thousands of glucose molecules covalently bonded together; See Fig 2.10c, page 49. 1. 2.

starch = plant storage carbohydrate; (24-30 glucose molecules); glycogen = animal storage carbohydrate; (12 glucose molecules); stored in liver and skeletal muscle.

*Polymers are broken down by hydrolysis resulting in monosaccharides. 4.

Function = energy storage! *How is the energy that is stored in carbohydrates released? *CELLULAR RESPIRATION OVERVIEW: oxygen glucose -------------------> H20 + CO2 energy (ATP) 29

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) C.

LIPIDS 1. 2.

contain C, H, and O, but much less O than in carbohydrates; types of lipids: a.

Fats

(See Fig 2.12, page 50):

m

monomers (building blocks) = triglycerides (glycerol + 3 fatty acids); saturated vs. unsaturated fats: See Fig 2.11, page 49.

m

m

1.

saturated fats: a. have only single bonds between the carbons in their fatty acid chains; b. are solid at room temperature; c. are animal fats; d. include bacon grease, lard, butter; e. are nutritionally "BAD" fat;

2.

unsaturated fats: a. have one or more double bond between the carbons in their fatty acid chains; b. are liquid at RT (oils); c. are plant fats; d. include corn and olive oil, e. are nutritionally "GOOD" fat;

Function = energy source.

b.

Phospholipids (See Fig 2.13, page 50) m triglyceride with the substitution of a phosphate group (PO4-) for one fatty acid chain; m Function = major cell membrane component.

c.

Steroids (See Fig 2.14, page 50) m four interconnected carbon rings; m Example is cholesterol; m Function = compose cell membranes; chemical messengers (hormones). 30

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) D.

PROTEINS 1.

Monomers = amino acids a.

Amino Acid Structure (See Fig 2.15a, page 51) asymmetric C HO H H3N ------- C ------- C

+

ROamino group side chain b.

Types of amino acids m m

2.

carboxyl group

20 different based on R-groups or side-chains See Fig 2.15b, page 51 for some examples.

Polymers are formed by dehydration synthesis between the amino group of one amino acid and the carboxyl group of a 2nd amino acid. See Fig 2.16a, page 52. a. b.

3.

Bond formed = a peptide bond Length of amino acid chains may vary: m peptide = 2-9 aa’s; m polypeptide = 10-100 aa’s; m protein = over 100 aa’s.

Protein Structure (See Fig 2.16-2.18 on page 52). a. Primary (1o) = sequence of amino acids; b. Secondary (2o) = twisting of amino acid chain; due to hydrogen bonding; c. Tertiary (3o) = folding of the amino acid chain; due to ionic bonds, disulfide bridges, and hydrophobic interactions; d. Quaternary (4o) = interactions between different amino acid chains; * A protein must be in its quaternary structure to be functional!!!! 31

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) D.

Proteins (continued) 4.

Denaturation of Proteins = the loss of 3 dimensional conformation of a protein; = loss of function. a.

Reasons for denaturation: m m m m

5.

extreme pH values; extreme temperature values; harsh chemicals (disrupt bonding); high salt concentrations.

Functions of Proteins a. b. c. d. e. f.

structure = keratin in hair, nails and skin; transport = hemoglobin; chemical messengers = hormones & neurotransmitters; movement = actin and myosin in muscle; defense = antibodies; catalysts = ENZYMES.

32

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) E.

NUCLEIC ACIDS 1.

Monomers = nucleotides; a.

Nucleotide structure = 3 parts: m m m

2.

See Fig 2.19, page 53.

pentose sugar (5-C); nitrogenous base; 1. purine (double ring) or 2. pyrimidine (single ring); phosphate group.

Polymers are formed by bonding between the sugar of one nucleotide and the phosphate group of a second nucleotide = sugar/phosphate backbone; See Fig 2.20, page 53.

3.

Types of Nucleic Acids a.

DEOXYRIBONUCLEIC ACID = DNA m

Structure 1. Sugar = deoxyribose; 2. Bases = adenine (A), thymine (T), cytosine(C), guanine (G); 3. double stranded (resembles ladder); strands held together by H-bonds between bases on opposite strands: a. b. 4.

A=T; C=G.

double helix (ladder is twisted).

m

Function = genetic material (i.e. genes, chromosomes).

m

DNA contains all necessary information needed to sustain and reproduce life!

33

CHAPTER 2: CHEMISTRY VI.

Organic Compounds (continued) E.

NUCLEIC ACIDS 3.

types of nucleic acids b.

Ribonucleic Acid = RNA m

Structure 1. 2. 3.

4.

Sugar = ribose; Bases = A,G,C, and uracil (not thymine) single stranded.

Function = transport DNA code during protein synthesis.

See Chart 2.8, page 53 for a summary of the four organic compounds in living things.

34

CHAPTER 2: CHEMICAL BASIS OF LIFE ORGANIC MOLECULE SUMMARY TABLE (key on page 38 of outline) Organic Molecule Composed of what atoms? Building Blocks (monomers)

Specific types & functions of monomers

Specific types and functions of polymers

Other

35

CHAPTER 2: CHEMICAL BASIS OF LIFE SUBATOMIC PARTICLE SUMMARY TABLE (outline page 17) SUBATOMIC PARTICLE

CHARGE

LOCATION

MASS (WEIGHT)

PROTON

POSITIVE

NUCLEUS

1 amu

NEUTRON

ZERO (NEUTRAL)

NUCLEUS

1 amu

ELECTRON

NEGATIVE

SHELLS OR ORBITALS AROUND NUCLEUS

0

CHEMICAL BOND SUMMARY TABLE (outline page 21) TYPE OF BOND

DEFINITION

DESCRIPTION

EXAMPLE

IONIC

when atoms lose or gain electrons becoming ions, and then oppositely charged ions are attracted to one another

bond is broken by water

salts, NaCl

COVALENT

when 1 or more pair(s) of electrons are shared by atoms

strong bond

the bonds holding H20 together (intra-), CO2

HYDROGEN

when a (slightly positive) hydrogen atom that is already covalently bonded to something else is attracted to a slightly negative atom.

Very weak bond; in molecules whose purpose is to easily break and then come back together

reactions between water molecules (inter-; i.e. ice to water to gas); DNA chains

36

CHAPTER 2: CHEMICAL BASIS OF LIFE Chemical Reaction Comparison Table (outline page 23)

SYNTHESIS REACTIONS

DEGRADATION RXN’S

GENERAL DESCRIPTION

Synthesis involves the building of a large molecule (polymer) from smaller building blocks (monomer).

Degradation involves the breakdown of polymer into individual monomers.

DESCRIPTIVE TERMS

building constructive anabolic

breakdown digestive decomposition catabolic

BOND FORMATION OR BREAKING?

Bonds are formed.

Bonds are broken.

IS ENERGY REQUIRED OR RELEASED? NAME THAT TERM.

Energy is required to form the bond. Endergonic

Energy is released when the bond is broken. Exergonic

HOW IS WATER INVOLVED? NAME THAT TERM.

Water is released when he bond is formed. Dehydration

Water is required to break the bond. Hydrolysis

EXAMPLE

Building a protein from individual amino acids; Building a triglyceride from glycerol and 3 fatty acids, etc

Breaking a protein into individual amino acids; Breaking starch down into monosaccharides, etc.

37

CHAPTER 2: CHEMICAL BASIS OF LIFE ORGANIC MOLECULE SUMMARY TABLE (outline page 32) Organic Molecule

Carbohydrates (sugars)

Lipids (Fats)

Proteins

Nucleic Acids

Composed of what atoms?

C, H, 0

C, H, O

C, H, O, N, S

C, H, O, N, P

Building Blocks (monomers)

Monosaccharides or hexoses

Triglycerides: glycerol and 3 fatty acids

amino acids

nucleotides: pentose sugar, phosphate, nitrogen base

Specific types & functions of monomers

glucose, fructose, galactose:

TG = energy

20 different amino acids

N/A

proteins (>100 amino acids); Many functions: ENZYMES, antibodies, struc-ture, transport, chemical messengers, storage

DNA = deoxyribo=nucl eic acid; genetic material; RNA= ribonucleic acid; aids DNA in protein synthesis.

energy

Phospholipid = cell membrane component Steroid = cell membrane component and chemical messenger (i.e. cholesterol)

Specific types and functions of polymers

Other

Disaccharides: sucrose, lactose, maltose; energy _____________ Polysaccharides : Starch (plant); Glycogen (animal); energy storage.

N/A

Saturated (only single bonds between C’s in fa chain) vs. Unsaturated (at least 1 double bond in fa chain)

DNA controls cellular activity by instructing our cells what proteins to make (i.e. Enzymes).

38

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